Oxidising Agent
An oxidising agent is a substance that oxidises another substance and, in doing so, is itself reduced — it accepts (gains) electrons from the species being oxidised. Because it gains electrons, the oxidation number of at least one element in the oxidising agent decreases during the reaction; this is the definitive chemical test for identifying an oxidising agent in any redox equation. Common oxidising agents encountered at IGCSE level include acidified potassium manganate(VII) (KMnO₄), oxygen gas (O₂), chlorine gas (Cl₂), hydrogen peroxide (H₂O₂), and the halogens in general during displacement reactions. It is essential to distinguish between the agent (the substance that does the oxidising) and the process of oxidation (the loss of electrons by the other substance); the oxidising agent is always the one gaining electrons, not losing them. A useful practical test for an oxidising agent is to add it to a solution of potassium iodide (KI): if the oxidising agent is present, the colourless iodide ions are oxidised to brown iodine (I₂), so the solution turns from colourless to brown. Many oxidising agents also produce characteristic colour changes when they themselves are reduced, such as purple acidified KMnO₄ turning colourless (Mn²⁺) or orange acidified K₂Cr₂O₇ turning green (Cr³⁺), which makes them useful in test-tube identification reactions and redox titrations.
Definition and Core Principle
The oxidising agent is fundamentally defined by what happens to it, not what it does to others: an oxidising agent is the species that becomes reduced. The term “oxidising” in its name refers to its effect on the other reactant — it causes the other substance to be oxidised — but the chemical change the agent itself undergoes is always reduction. This is a critical distinction.
The process can be remembered through the mnemonic “OIL RIG”:
- Oxidation Is Loss (of electrons)
- Reduction Is Gain (of electrons)
From the perspective of the oxidising agent:
- The oxidising agent gains electrons → it is reduced
- It causes the other substance to lose electrons → the other substance is oxidised
In every redox reaction, there are always two agents at work simultaneously: an oxidising agent and a reducing agent. One cannot exist without the other, because electrons transferred by one species must be accepted by another.
How to Identify an Oxidising Agent
To identify the oxidising agent in a redox reaction, follow these steps:
- Assign oxidation numbers to all atoms in both the reactants and products.
- Identify which elements change oxidation number.
- Find the element (or elements) whose oxidation number decreases — this means it has gained electrons and been reduced.
- The substance containing that element is the oxidising agent.
Worked example: The thermite reaction
Fe₂O₃ + 2Al → 2Fe + Al₂O₃
- In Fe₂O₃, iron has an oxidation number of +3. In the product, Fe is a free element, so its oxidation number is 0. Iron’s oxidation number decreases from +3 to 0: iron has gained 3 electrons — it has been reduced.
- In Al, the oxidation number is 0 (free element). In Al₂O₃, aluminium has an oxidation number of +3. Aluminium’s oxidation number increases from 0 to +3: aluminium has lost 3 electrons — it has been oxidised.
- The substance that contains the element being reduced is Fe₂O₃. Therefore, Fe₂O₃ (iron(III) oxide) is the oxidising agent.
- The substance that contains the element being oxidised is Al. Therefore, Al (aluminium) is the reducing agent.
Worked example: Displacement of iodine by chlorine
Cl₂ + 2KI → 2KCl + I₂
Ionic equation: Cl₂ + 2I⁻ → 2Cl⁻ + I₂
- Chlorine: oxidation number goes from 0 (Cl₂) to -1 (Cl⁻). Decrease of 1 → chlorine is reduced. Cl₂ is the oxidising agent.
- Iodine (in I⁻): oxidation number goes from -1 to 0 (I₂). Increase of 1 → iodide is oxidised. I⁻ (or KI) is the reducing agent.
Common Oxidising Agents at IGCSE
| Oxidising Agent | Formula | What It Does | Colour Change Observed | Key Half-Equation |
|---|---|---|---|---|
| Acidified potassium manganate(VII) | KMnO₄ (with H₂SO₄) | Oxidises a wide range of reducing agents; used as a test for reducing agents and in redox titrations | Purple → colourless (Mn²⁺ formed) | MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O |
| Oxygen | O₂ | Supports combustion; oxidises metals to form metal oxides, non-metals to form non-metal oxides; rusts iron | Not directly visible (O₂ is colourless, but the product e.g. rust is visible) | O₂ + 4e⁻ → 2O²⁻ |
| Chlorine | Cl₂ | Displaces less reactive halogens from their salts; used as a bleaching agent and in water purification | Pale green-yellow Cl₂ gas → colourless Cl⁻ solution (if testing on KI paper, turns it brown) | Cl₂ + 2e⁻ → 2Cl⁻ |
| Hydrogen peroxide | H₂O₂ | Oxidising agent in many reactions; decomposes to give oxygen (where it also acts as a reducing agent — it can disproportionate) | Decomposes to water and oxygen; colour change depends on the reaction context | H₂O₂ + 2H⁺ + 2e⁻ → 2H₂O |
| Acidified potassium dichromate(VI) | K₂Cr₂O₇ (with H₂SO₄) | Oxidises alcohols to carboxylic acids; used as a test for reducing agents such as ethanol or SO₂ | Orange → green (Cr³⁺ formed) | Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O |
| Bromine | Br₂ | Displaces iodide ions from solution; less powerful oxidising agent than chlorine | Orange-brown Br₂ → colourless Br⁻ | Br₂ + 2e⁻ → 2Br⁻ |
| Iodine | I₂ | Weakest halogen oxidising agent; can still oxidise some species | Brown I₂ → colourless I⁻ | I₂ + 2e⁻ → 2I⁻ |
Halogens as Oxidising Agents
The halogens act as oxidising agents, and their oxidising power decreases down Group 7:
Cl₂ > Br₂ > I₂ (in terms of oxidising ability)
This trend explains displacement reactions:
- Chlorine displaces both bromide and iodide ions because Cl₂ is a stronger oxidising agent than Br₂ or I₂.
- Bromine displaces iodide ions because Br₂ is a stronger oxidising agent than I₂, but cannot displace chloride ions.
- Iodine cannot displace either chloride or bromide ions because it is the weakest oxidising agent of the three.
Example displacement reaction:
Cl₂ (aq) + 2KBr (aq) → 2KCl (aq) + Br₂ (aq)
Ionic: Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂
Here, Cl₂ is the oxidising agent (it gains electrons and is reduced to Cl⁻), and Br⁻ is the reducing agent (it loses electrons and is oxidised to Br₂). The orange-brown colour of bromine appears in the solution.
Oxidising Agent vs Oxidation
This distinction is a frequent source of confusion and a common exam pitfall:
| Concept | Definition | What Happens to Electrons | Oxidation Number Change | Example |
|---|---|---|---|---|
| Oxidation | The process of losing electrons | Electrons are lost by the substance being oxidised | Oxidation number increases | Fe²⁺ → Fe³⁺ + e⁻ (iron has been oxidised) |
| Oxidising Agent | The substance that causes oxidation by accepting electrons | Electrons are gained by the oxidising agent (it is reduced) | Oxidation number of an element within it decreases | MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O (MnO₄⁻ is the oxidising agent) |
| Reduction | The process of gaining electrons | Electrons are gained by the substance being reduced | Oxidation number decreases | MnO₄⁻ gains electrons and is reduced |
| Reducing Agent | The substance that causes reduction by donating electrons | Electrons are lost by the reducing agent (it is oxidised) | Oxidation number of an element within it increases | Fe²⁺ is the reducing agent (it donated electrons to MnO₄⁻) |
Memory aid: The oxidising agent is the one that gets reduced. The reducing agent is the one that gets oxidised. This sounds contradictory at first, but it makes sense: the agent does something to the other substance and in the process suffers the opposite fate itself.
Sources, Common Misconceptions
Sources
- Cambridge IGCSE Chemistry 0620 Syllabus, Topic 6: Chemical Reactions (Redox)
- Edexcel IGCSE Chemistry Specification, Section 2: Inorganic Chemistry (Reactivity Series)
- Roger Norris, Cambridge IGCSE Chemistry Revision Guide
- Lawrie Ryan and Roger Norris, Cambridge IGCSE & O Level Chemistry
Common Misconceptions
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“The oxidising agent is the substance that gets oxidised.” This is completely wrong and is the single most common error students make. The oxidising agent is itself reduced (gains electrons). It causes the other substance to be oxidised. Think: an “oxidising agent” is an agent of oxidation — it brings about oxidation in something else, while it itself undergoes reduction.
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“Oxidation and the oxidising agent are the same thing.” They are not. Oxidation is a process (the loss of electrons), while the oxidising agent is a substance (the chemical species that accepts those electrons). This is a bit like confusing the act of eating with the food itself.
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“Only oxygen can be an oxidising agent.” This misconception comes from the historical definition of oxidation as “combination with oxygen.” While oxygen is a very important oxidising agent, many other substances can accept electrons and act as oxidising agents — including halogens, KMnO₄, K₂Cr₂O₇, and even metal ions such as Cu²⁺ in displacement reactions.
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“If the oxidation number goes up, the substance is the oxidising agent.” Exactly the opposite. If the oxidation number of an element increases, that element has lost electrons — it has been oxidised — meaning it is acting as a reducing agent. The oxidising agent is the one whose oxidation number decreases.
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“An oxidising agent and a reducing agent are the same substance.” A substance can sometimes act as an oxidising agent in one reaction and a reducing agent in another (depending on the conditions and what it reacts with), but it cannot be both at the same time in a single reaction. Key exceptions include disproportionation reactions, such as the decomposition of hydrogen peroxide (2H₂O₂ → 2H₂O + O₂), where the same species is simultaneously oxidised and reduced — but in this case H₂O₂ is not acting as an “agent” towards something else; it is reacting with itself.